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9 oxygenfluorineneon


Name, symbol, number fluorine, F, 9
Chemical serieshalogens
Group, period, block 172, p
AppearanceYellowish brown gas
Standard atomic weight 18.9984032(5) g·mol−1
Electron configuration 1s2 2s2 2p5
Electrons per shell 2, 7
Physical properties
Density(0 °C, 101.325 kPa)
1.7 g/L
Melting point53.53 K
(−219.62 °C, −363.32 °F)
Boiling point85.03 K
(−188.12 °C, −306.62 °F)
Critical point144.13 K, 5.172 MPa
Heat of fusion(F2) 0.510 kJ·mol−1
Heat of vaporization(F2) 6.62 kJ·mol−1
Heat capacity(25 °C) (F2)
31.304 J·mol−1·K−1
Vapor pressure
P/Pa 1 10 100 1 k 10 k 100 k
at T/K 38 44 50 58 69 85
Atomic properties
Crystal structurecubic
Oxidation states−1
(strongly acidic oxide)
Electronegativity3.98 (Pauling scale)
Ionization energies
1st: 1681.0 kJ·mol−1
2nd: 3374.2 kJ·mol−1
3rd: 6050.4 kJ·mol−1
Atomic radius50 pm
Atomic radius (calc.)42 pm
Covalent radius71 pm
(see covalent radius of fluorine)
Van der Waals radius147 pm
Magnetic orderingnonmagnetic
Thermal conductivity(300 K) 27.7 m W·m−1·K−1
CAS registry number7782-41-4
Selected isotopes
Main article: Isotopes of fluorine
iso NA half-life DM DE (MeV) DP
18F syn 109.77 min ε 1.656 18O
19F 100% F is stable with 10 neutrons
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Fluorine (pronounced /ˈflʊəriːn/, Latin: fluere, meaning "to flow"), is the chemical element with the symbol F and atomic number 9. Atomic fluorine is univalent and is the most chemically reactive and electronegative of all the elements. In its elementally isolated (pure) form, fluorine is a poisonous, pale, yellowish brown gas, with chemical formula F2. Like other halogens, molecular fluorine is highly dangerous; it causes severe chemical burns on contact with skin.

Fluorine's large electronegativity and small atomic radius gives it interesting bonding characteristics, particularly in conjunction with carbon. See covalent radius of fluorine.


Notable characteristics

Pure fluorine (F2) is a corrosive pale yellow or brown[1] gas that is a powerful oxidizing agent. It is the most reactive and most electronegative of all the elements (4.0), and readily forms compounds with most other elements. It has an oxidation number -1, except when bonded to another fluorine in F2 which gives it an oxidation number of 0. Fluorine even combines with argon, krypton, xenon, and radon. Even in dark, cool conditions, fluorine reacts explosively with hydrogen. It is so reactive that metals, and even water, as well as other substances, burn with a bright flame in a jet of fluorine gas. It is far too reactive to be found in elemental form. In moist air it reacts with water to form also-dangerous hydrofluoric acid.

In aqueous solution, fluorine commonly occurs as the fluoride ion F, although highly diluted HF is such a weak acid that substantial amounts of it are present in any water solution of fluoride at near neutral pH. Other forms are fluoro-complexes, such as [FeF4], or H2F+.

Fluorides are compounds that combine fluorine with some positively charged counterpart. They often consist of crystalline ionic salts. Fluorine compounds with metals are among the most stable of salts.


Chemical uses:

  • Atomic fluorine and molecular fluorine are used for plasma etching in semiconductor manufacturing, flat panel display production and MEMS (microelectromechanical systems) fabrication[2]. Xenon difluoride is also used for this last purpose.
  • Hydrofluoric acid (chemical formula HF) is used to etch glass in light bulbs and other products.
  • Fluorine is indirectly used in the production of low friction plastics such as Teflon, and in halons such as freon.
  • Along with some of its compounds, fluorine is used in the production of pure uranium from uranium hexafluoride and in the synthesis of numerous commercial fluorochemicals, including vitally important pharmaceuticals, agrochemical compounds, lubricants, and textiles.
  • Fluorochlorohydrocarbons are used extensively in air conditioning and in refrigeration. Chlorofluorocarbons have been banned for these applications because they contribute to ozone destruction and the ozone hole. Interestingly, since it is chlorine and bromine radicals which harm the ozone layer, not fluorine, compounds which do not contain chlorine or bromine but contain only fluorine, carbon and hydrogen (called hydrofluorocarbons) are not on the EPA list of ozone-depleting substances,[3] and have been widely used as replacements for the chlorine- and bromine-containing fluorocarbons. Hydrofluorocarbons do have a greenhouse effect, but a small one compared with carbon dioxide and methane.
  • Sulfur hexafluoride is an extremely inert and nontoxic gas, very useful as an insulator in high-voltage electrical equipment. It does not occur in nature, so it is a useful tracer gas, though as an exceptionally potent greenhouse gas its use in unenclosed systems is inadvisable.
  • Sodium hexafluoroaluminate (cryolite), is used in the electrolysis of aluminium.
  • In much higher concentrations, sodium fluoride has been used as an insecticide, especially against cockroaches.
  • Fluorides have been used in the past to help molten metal flow, hence the name.
  • Some researchers including US space scientists in the early 1960s have studied elemental fluorine gas as a possible rocket propellant due to its exceptionally high specific impulse. The experiments failed because fluorine proved difficult to handle, and its combustion products proved extremely toxic and corrosive.
  • Polytetrafluoroethylene, also known as the non-stick Teflon surface in baking pans.
  • Compounds of fluorine such as fluoropolymers, potassium fluoride and cryolite are utilized in applications such as anti-reflective coatings and dichroic mirrors on account of their unusually low refractive index.

Dental and medical uses:


Fluorine forms a variety of very different compounds, owing to its small atomic size and covalent behavior, and on the other hand, its oxidizing ability and extreme electronegativity. For example, hydrofluoric acid is extremely dangerous, while in synthetic drugs incorporating an aromatic ring (e.g. flumazenil), fluorine is used to prevent toxication or to delay metabolism.

The fluoride ion is basic, therefore hydrofluoric acid is a weak acid in water solution. However, water is not an inert solvent in this case: when less basic solvents such as anhydrous acetic acid are used, hydrofluoric acid is the strongest of the hydrohalogenic acids. Also, owing to the basicity of the fluoride ion, soluble fluorides give basic water solutions. The fluoride ion is a Lewis base, and has a high affinity to certain elements such as calcium and silicon. For example, deprotection of silicon protecting groups is achieved with a fluoride. The fluoride ion is poisonous.

Fluorine as a freely reacting oxidant gives the strongest oxidants known. Chlorine trifluoride, for example, can burn water and sand, both compounds of a weaker oxidant, oxygen.

Fluorine compounds involving noble gases were first synthesised by Neil Bartlett in 1962—xenon hexafluoroplatinate, XePtF6, being the first. Fluorides of krypton and radon have also been prepared. Also argon fluorohydride has been prepared, although it is only stable at cryogenic temperatures.

The carbon-fluoride bond is covalent and very stable. The use of a fluorocarbon polymer, poly(tetrafluoroethene) or Teflon, is an example: it is thermostable and waterproof enough to be used in frying pans. Organofluorines may be safely used in applications such as drugs, without the risk of release of toxic fluoride. In synthetic drugs, toxication can be prevented. For example, an aromatic ring is useful but presents a safety problem: enzymes in the body metabolize some of them into poisonous epoxides. When the para position is substituted with fluorine, the aromatic ring is protected and epoxide is no longer produced.

Fluorine can often be substituted for hydrogen when it occurs in organic compounds. Through this mechanism, fluorine can have a very large number of compounds.


This element is recovered from fluorite, cryolite, and fluorapatite.

For a list of fluorine compounds, see here.


Fluorine in the form of fluorspar (also called fluorite, calcium fluoride) was described in 1530 by Georgius Agricola for its use as a flux, [5] which is a substance that is used to promote the fusion of metals or minerals. In 1670 Schwanhard found that glass was etched when it was exposed to fluorspar that was treated with acid. Carl Wilhelm Scheele and many later researchers, including Humphry Davy, Caroline Menard, Gay-Lussac, Antoine Lavoisier, and Louis Thenard all would experiment with hydrofluoric acid, easily obtained by treating calcium fluoride (fluorspar) with concentrated sulfuric acid.

It was eventually realized that hydrofluoric acid contained a previously unknown element. This element was not isolated for many years after this, due to its extreme reactivity; fluorine can only be prepared from its compounds electrolytically, and then it immediately attacks any susceptible materials in the area. Finally, in 1886, elemental fluorine was isolated by Henri Moissan after almost 74 years of continuous effort by other chemists.[6] It was an effort which cost several researchers their health or even their lives. The derivation of elemental fluorine from hydrofluoric acid is exceptionally dangerous, killing or blinding several scientists who attempted early experiments on this halogen. These men came to be referred to as "fluorine martyrs". For Moissan, it earned him the 1906 Nobel Prize in chemistry (Moissan himself lived to be 54, and it is not clear whether his fluorine work shortened his life).

The first large-scale production of fluorine was needed for the atomic bomb Manhattan project in World War II where the compound uranium hexafluoride (UF6) was needed as a gaseous carrier of uranium to separate the 235U and 238U isotopes of uranium. Today both the gaseous diffusion process and the gas centrifuge process use gaseous UF6 to produce enriched uranium for nuclear power applications. In the Manhattan Project, it was found that elemental fluorine was present whenever UF6 was, due to the spontaneous decomposition of this compound into UF4 and F2. The corrosion problem due to the F2 was eventually solved by electrolytically coating all UF6 carrying piping with nickel metal, which resists fluorine's attack. Joints and flexible parts were made from teflon, then a very recently discovered fluorocarbon plastic which was not attacked by F2.


Elemental fluorine is prepared industrially by Moissan's original process: electrolysis of anhydrous HF in which KHF2 has been dissolved to provide enough ions for conduction to take place.

In 1986, when preparing for a conference to celebrate the 100th anniversary of the discovery of fluorine, Karl Christe discovered a purely chemical preparation by reacting together at 150 °C solutions in anhydrous HF of K2MnF6 and of SbF5. The reaction is:

K2MnF6 + 2SbF5 → 2KSbF6 + MnF3 + ½F2

This is not a practical synthesis, but demonstrates that electrolysis is not essential.


Main article: fluoride poisoning

Both elemental fluorine and fluoride ions are highly toxic and must be handled with great care and any contact with skin and eyes should be strictly avoided. When it is a free element, fluorine has a characteristic pungent odor that is detectable in concentrations as low as 20 nL/L. All equipment must be passivated before exposure to fluorine.

Contact of exposed skin with hydrofluoric acid solutions poses one of the most extreme and insidious industrial threats—one which is exacerbated by the fact that hydrofluoric acid damages nerves in such a way as to make such burns initially painless. The hydrofluoric acid molecule is capable of rapidly migrating through lipid layers of cells which would ordinarily stop an ionized acid, and the burns are typically deep. HF may react with calcium, permanently damaging the bone. More seriously, reaction with the body's calcium can cause cardiac arrhythmias, followed by cardiac arrest brought on by sudden chemical changes within the body. These cannot always be prevented with local or intravenous injection of calcium salts. Hydrofluoric acid spills over just 2.5% of the body's surface area (about 75 in2 or 5 dm2), despite copious immediate washing, have been fatal.[7] If the patient survives, hydrofluoric acid burns typically produce open wounds of an especially slow-healing nature.

Elemental fluorine is a powerful oxidizer which can cause organic material, combustibles, or other flammable materials to ignite.

Fluorocarbons are generally inert and nontoxic; the electronegativity of fluorine means that a nearby fluorine atom makes a carboxylic acid group very much more reactive. For example, trifluoroacetic acid is 100,000 times stronger than acetic acid.

See also


  • Los Alamos National Laboratory – Fluorine
  1. ^ Theodore Gray. Real visible fluorine. The Wooden Periodic Table.
  2. ^ Leonel R Arana, Nuria de Mas, Raymond Schmidt, Aleksander J Franz, Martin A Schmidt and Klavs F Jensen, Isotropic etching of silicon in fluorine gas for MEMS micromachining , J. Micromech. Microeng. 17 , 2007, pp. 384-392.
  3. ^ Class I Ozone-Depleting Substances. Ozone Depletion. U.S. Environmental Protection Agency.
  4. ^
  5. ^ Fluoride History Discovery of fluorine
  6. ^ H. Moissan (1886). "Action d'un courant électrique sur l'acide fluorhydrique anhydre". Comptes rendus hebdomadaires des séances de l'Académie des sciences 102: 1543-1544.
  7. ^ [1]
This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Fluorine". A list of authors is available in Wikipedia.
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